Revision notes for AQA GCSE Chemistry Collision theory and activation energy. Open the guide for explanations and worked examples. Written against the AQA GCSE Chemistry (8462) specification, so the content matches what's examinable rather than general Chemistry background.
Collision theory and activation energy
What you'll learn
- Why particles must collide before a chemical reaction can happen.
- What activation energy means and how it links to reaction profiles.
- How concentration, gas pressure, surface area and temperature affect rate.
- How to explain these changes using collision theory, not just memorise them.
Starting point: what does “rate of reaction” mean?
The rate of reaction tells you how fast reactants are used up or products are made. A fast reaction makes product quickly; a slow reaction makes product gradually.
Definition
Rate of reaction
The rate of reaction is the change in the amount of reactant or product in a given time.
A common mean rate equation is:
mean rate of reaction=amount of reactant used or product formedtime\text{mean rate of reaction} = \frac{\text{amount of reactant used or product formed}}{\text{time}}mean rate of reaction=timeamount of reactant used or product formedThe units depend on what you measure: for example, grams per second (g/s), cubic centimetres per second (cm³/s), or moles per second (mol/s).
Example
Calculating a mean rate
A reaction produces 80 cm³ of gas in 40 seconds.
- Choose the correct rate equation: mean rate is product formed divided by time.
- Substitute the values: 80 cm340 s=2.0 cm3/s\frac{80\ \text{cm}^3}{40\ \text{s}} = 2.0\ \text{cm}^3/\text{s}40 s80 cm3=2.0 cm3/s.
- The mean rate is 2.0 cm³/s, meaning the reaction produced 2.0 cm³ of gas each second on average.
Collision theory: the particle explanation
Chemical reactions happen between particles: atoms, ions or molecules. These particles are moving around, especially in liquids and gases.
Definition
Collision theory
Collision theory explains reaction rates by saying that reacting particles must collide with each other for a reaction to happen.
But not every collision causes a reaction. Some particles collide and simply bounce apart.
Definition
Successful collision
A successful collision is a collision that leads to a chemical reaction and forms products.
Key Idea
The big idea
The faster the rate of reaction, the more successful collisions happen each second.
So, when an exam asks you to explain why a reaction is faster, your answer usually needs this chain:
- Particles collide more often and/or with more energy.
- More collisions are successful.
- The rate of reaction increases.
Activation energy
Even if particles collide, they need enough energy to react. This energy is needed to start breaking bonds and allow new bonds to form.
Definition
Activation energy
The activation energy is the minimum amount of energy that reacting particles must have when they collide for a reaction to occur.
A reaction profile diagram shows energy during a reaction. The activation energy is measured from the energy level of the reactants up to the highest point of the curve.
Example
Reading activation energy from a profile
A reaction profile shows reactants at 30 kJ/mol and the highest point of the curve at 95 kJ/mol.
- Activation energy is the energy gap from the reactants to the peak.
- Calculate the difference: 95 kJ/mol−30 kJ/mol=65 kJ/mol95\ \text{kJ/mol} - 30\ \text{kJ/mol} = 65\ \text{kJ/mol}95 kJ/mol−30 kJ/mol=65 kJ/mol.
- The activation energy is 65 kJ/mol.
Common Mistake
Collisions are not enough
Do not just write “more collisions”. To explain a faster reaction properly, say more successful collisions per second.
Increasing concentration in solution
Definition
Concentration
Concentration is the amount of solute dissolved in a certain volume of solution. At GCSE it is often measured in moles per cubic decimetre, written as mol/dm³.
If a solution is more concentrated, there are more reactant particles in the same volume. The particles are closer together, so they collide more frequently.
For example, hydrochloric acid reacts with magnesium:
Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
If the acid is more concentrated, acid particles hit the magnesium surface more often. This gives more successful collisions each second, so the rate increases.
Example
Comparing acid concentrations
Two experiments use the same mass of magnesium ribbon, the same temperature and the same volume of hydrochloric acid. Experiment A uses 0.5 mol/dm³ acid. Experiment B uses 1.0 mol/dm³ acid.
- The only important change is the acid concentration, because the mass of magnesium, temperature and volume are controlled.
- Experiment B has more acid particles in the same volume, so acid particles collide with the magnesium surface more frequently.
- More frequent collisions means more successful collisions per second.
- Experiment B has the faster rate of reaction.
Common Mistake
Be careful with “doubles”
If the concentration doubles, the collision frequency often increases roughly in proportion. However, only say the rate doubles if the question or data shows direct proportionality.
Increasing pressure for reacting gases
Definition
Pressure
For a gas, pressure is caused by gas particles colliding with the walls of their container.
Increasing the pressure of reacting gases means the same number of gas particles is squeezed into a smaller volume, or more gas particles are added to the same volume.
Either way, the particles are closer together. They collide more frequently, so there are more successful collisions per second and the rate of reaction increases.
This is very similar to increasing concentration, but it applies to gases rather than solutions.
Tip
Concentration and pressure link
Higher concentration in a solution and higher pressure in a gas both mean: more particles in a given volume, so more frequent collisions.
Increasing surface area of a solid
Many reactions involve a solid reacting with a solution or gas. Only the particles on the surface of the solid can collide with the other reactant.
If you break a solid into smaller pieces, the total amount of solid may stay the same, but more of it is exposed.
Definition
Surface area to volume ratio
The surface area to volume ratio compares how much exposed surface a solid has compared with its volume.
A powder has a much larger surface area to volume ratio than one large lump. This means more solid particles are available to collide at the same time, so the reaction is faster.
Example
Calculating surface area to volume ratio
A solid cube has side length 2 cm. It is cut into eight smaller cubes, each with side length 1 cm.
- For the original cube, calculate surface area and volume: 6s2=6×(2 cm)2=24 cm26s^2 = 6 \times (2\ \text{cm})^2 = 24\ \text{cm}^26s2=6×(2 cm)2=24 cm2, and s3=(2 cm)3=8 cm3s^3 = (2\ \text{cm})^3 = 8\ \text{cm}^3s3=(2 cm)3=8 cm3.
- For the eight smaller cubes, calculate total surface area: 8×6×(1 cm)2=48 cm28 \times 6 \times (1\ \text{cm})^2 = 48\ \text{cm}^28×6×(1 cm)2=48 cm2. The total volume is still 8 cm³.
- Compare the ratios: the original cube has 24 cm² of surface for 8 cm³ of solid, but the smaller cubes have 48 cm² of surface for the same volume.
- Cutting the cube doubles the exposed surface area, so collisions at the surface can happen more frequently and the reaction is faster.
Common Mistake
Same mass, different rate
Breaking a solid into smaller pieces does not create more solid. It increases the exposed surface area, so the same amount of solid can react faster.
Increasing temperature
Definition
Temperature
Temperature is a measure of the average kinetic energy of particles. Kinetic energy is the energy particles have because they are moving.
When you increase the temperature, two important things happen:
- Particles move faster, so they collide more frequently.
- Particles have more energy, so a larger proportion of collisions have energy equal to or greater than the activation energy.
The second point is especially important. A small temperature increase can cause a big increase in rate because many more particles now have enough energy to react.
Example
Explaining a temperature increase
A reaction is repeated at 20 °C and 40 °C, with all other conditions kept the same.
- At 40 °C, the reacting particles have greater average kinetic energy than at 20 °C.
- The particles move faster, so they collide more often.
- More importantly, a greater proportion of collisions have energy at least equal to the activation energy.
- Therefore, there are more successful collisions per second, so the reaction is faster at 40 °C.
Common Mistake
Temperature does not lower activation energy
Increasing temperature gives particles more energy. It does not lower the activation energy.
The GCSE explanation pattern
For this topic, exam questions often ask you to predict and explain what happens when a condition changes.
A strong answer usually follows this pattern:
- State what happens to the particles.
- Link this to collision frequency and/or collision energy.
- State what happens to the number of successful collisions per second.
- State what happens to the rate.
Key Idea
Factor summary
Increasing concentration, gas pressure or surface area increases the frequency of collisions. Increasing temperature increases collision frequency and makes collisions more energetic.
Exam technique
In the exam
- Use the phrase successful collisions per second when explaining rate changes.
- For temperature, mention both faster movement and more particles having at least the activation energy.
- For surface area, say that smaller pieces expose more surface particles; do not just say “the solid is smaller”.
Self review
Check yourself
- Why can particles collide without reacting?
- How does increasing gas pressure increase the rate of reaction?
- Why does powdered calcium carbonate react faster with acid than large marble chips of the same mass?
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